REVIEWTOPIC 5/15: THERMOCHEMISTRY / ENERGETICS
A. KEY DEFINITIONS (MEMORIZE!)
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In an exothermic reaction, energy is released to the surroundings.
In an endothermic reaction, energy is absorbed from the surroundings.
Standard enthalpy change of reaction is the heat energy transferred when molar amounts of substances in a
balanced equation react under standard thermodynamic conditions
Average bond enthalpy is the energy (in kJ) needed to break one mole of a particular covalent bond in the gaseous
state into gaseous atoms. It is an average because it is averaged over all the environments that the bond is found
in.
Standard conditions in thermodynamics are a temperature of 298 K (25°C), a pressure of 100 kPa (1.01 x105 Pa or
1 atmosphere), and 1 molar concentration of solutions. (**Don’t confuse with: Standard temperature and
pressure (STP / gases) – 273 K (0°C) and 101.3 kPa….molar volume is 22.4 dm3.)
The standard enthalpy change of formation, ∆Hf°, is the energy change when one mole of a substance is formed
from its elements under standard conditions. (e.g., Na(s) + ½Cl2(g)  NaCl(s) ∆Hf° = -411 kJ mol-1)
The standard enthalpy change of combustion, ∆Hc°, is the energy change when one mole of a substance burns
completely in oxygen.
Lattice enthalpy is the energy change associated when the ionic lattice is broken down to produce gaseous ions
[MX(s) + energy  M+(g) + X-(g)]] *The breaking down of an ionic lattice is endothermic.
Electron affinity is the energy change when a gaseous atom gains an electron:[X(g) + e-  X-(g)]
OTHERS: enthalpy of solution, hydration enthalpy
B. YOU SHOULD KNOW:
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That combustion (burning in oxygen) and neutralization are exothermic processes
∆H (the enthalpy change) is negative for exothermic reactions and ∆H is positive for endothermic reactions
∆H =  [n∆Hf(products)] -  [n∆Hf(reactants)], which means that the enthalpy change is the sum of the enthalpies
of formation of the products minus the sum of the enthalpies of formation of the reactants.
∆H =  [n∆Hc(reactants)] -  [n∆Hc(products)], which means that the enthalpy change is the sum of the
enthalpies of combustion of the reactants minus the enthalpies of combustion of the products.
q
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nsystem
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the equation q = mcΔT, and use it to calculate molar enthalpy changes (∆H =
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in experiments measuring enthalpy changes, the greatest error is always heat loss
Born-Haber cycles apply Hess’s law to ionic compounds and are often used to calculate lattice enthalpies
Both the charge and the size of ions affect the lattice energy of an ionic compound
Lattice enthalpies increase with higher ionic charges and smaller ionic radii
Differences between theoretical and experimental values for lattice enthalpies suggest that an ionic compound has
got some covalent character
Entropy is a measure of the disorder in a system and is represented using the symbol S
Entropy increases when substances are heated and when amounts of substances increase
Entropy is greatest in gases and least in solids
Entropy change, ∆S, is positive when entropy increases from reactants to products and negative when entropy
decreases
Gibbs free energy change, ∆G, tells us about the spontaneity of a reaction – if it is negative, a reaction will be
spontaneous (can happen without assistance at that temperature).
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C. YOU SHOULD BE ABLE TO:
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Recognize an exothermic reaction as one that raises the temperature of the surroundings
Recognize an endothermic reaction as one that leads to a decrease in temperature of the surroundings
Use enthalpy level diagrams
Understand that in an exothermic reaction the products are of lower potential energy than the reactants
Understand that in an endothermic reaction the reactants are of lower potential energy than the products
Explain that bond breaking uses energy (is endothermic) and bond making releases energy (is exothermic)
Discuss the bond strength in ozone relative to oxygen in its importance to the atmosphere
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Appreciate that it is the balance between bond breaking and bond making that determines whether a reaction will
be exothermic or endothermic
Work out the enthalpy change in a reaction using enthalpy changes of formation and combustion
Design experiments to measure the enthalpy change of combustion using the energy evolved to heat a known
mass of water
Design experiments to measure the enthalpy change in aqueous solution
Calculate both empirical and molecular formulas from given data
Work out the enthalpy change in a reaction using known enthalpy changes of two or three reactions with enthalpy
cycles of rearranging equations (e.g., applying Hess’s Law)
Construct Born-Haber cycles for the oxides and chlorides of elements in group 1 and 2 of the periodic table.
Calculate enthalpy changes for Born-Haber or dissolution energy cycles
Calculate ∆S, using standard entropy values and using the equation: ∆S =  [nS(products)] -  [nS(reactants)]
Use the equation ∆G = ∆H - T∆S
Predict how temperature will affect the spontaneity of a reaction
D. BE PREPARED:
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Standard conditions in thermodynamics are a temperature of 298 K (25°C), a pressure of 101.3 kPa (1.01 x10 5 Pa
or 1 atmosphere), and 1 molar concentration of solutions.
The specific heat capacity (c) of a substance is the amount of energy needed to raise the temperature of 1 g of it by
1 K (1°C). The specific heat capacity of water and aqueous solutions is 4.18 Jg-1K-1.
Lattice enthalpies at 298 K are found in the data booklet and are endothermic.
When calculating ∆G, remember that temperature must be measured in Kelvin (not in Celsius). Also, in data tables,
∆H is given in kJ and ∆S is given in J, so ∆S needs to be divided by 1000 to convert it to kJ.
Be able to calculate ∆G using: ∆G =  [n∆Gf(products)] -  [n∆Gf(reactants)]
MOR PRACTICE QUESTIONS (TOPIC 5/15)
1. Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of
hydrogenation of ethene to form ethane. C2H4(g) + H2(g)  C2H6(g)
John used the average bond enthalpies from Table 10. Marit used the values of enthalpies of combustion from Table 12.
(a) Calculate the value for the enthalpy of hydrogenation of ethene obtained using the average bond enthalpies given in
Table 10. (2)
(b) Determine the value of the enthalpy of hydrogenation of ethene using the values for the enthalpies of combustion of
ethene, hydrogen, and ethane given in Table 12. (2)
(c) Suggest one reason why John’s answer is slightly less accurate than Marit’s answer and calculate the percent difference.
(d) John then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane.
C6H10(l) + H2(g)  C6H12(l)
(i)
Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene. (1)
(ii)
The percentage difference between the two methods (average bond enthalpies and enthalpies of
combustion) is greater for cyclohexene than it was for ethane. John’s hypothesis was that it would be the
same. Determine why the use of average bond enthalpies is less accurate for the cyclohexene equation
shown above, than it was for ethane. Deduce what extra information is needed to provide a more
accurate answer. (2)
2. Consider the following Born-Haber cycle. (The magnitude for each of the
enthalpy changes (a to e) are given in kJ mol-1 but their signs (+ or -) have
been omitted.
a) State the names for the enthalpy changes c and d.
b) Deduce which two of the enthalpy changes a to e have negative signs.
c) Determine the value for the enthalpy of formation of potassium bromide.
d) Explain why the quantitative value for the lattice enthalpy of calcium
bromide is larger than the value for the lattice enthalpy of potassium
bromide.